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Aqueous

Aqueous. Reactions. Reactions in aqueous solutions. Aqueous solution: Solution in which water is the solvent (dissolving agent). 3 major types of chemical processes of aqueous solutions: Precipitation reactions Acid-base reactions Redox reactions. Solution:

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Aqueous

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  1. Aqueous Reactions

  2. Reactions in aqueous solutions • Aqueous solution: • Solution in which water is the solvent (dissolving agent). • 3 major types of chemical processes of aqueous solutions: • Precipitation reactions • Acid-base reactions • Redox reactions

  3. Solution: • Homogenous mixture of 2 or more substances. • Solvent: • Dissolving medium, usually present in greater quantity. • Solute: • The other substance(s) in the solution.

  4. Electrolytic Properties • Electrolyte: • A substance whose aqueous solution forms ions; conducts electricity. • Ionic compounds. • Nonelectrolyte: • Substance that does not form ions in an aqueous solution; poor conductor. • Molecular compounds.

  5. Ionic compounds in water: • Dissociate into its component ions.

  6. Do not get to wrapped up in the difference between the terms ionization and dissociation. Consider them to mean the same thing, the separation of a substances ions. Equations showing ionization or dissociation

  7. Molecular compounds in H2O • Molecular compounds – nonmetal + nonmetal • Structural integrity of molecule is usually maintained meaning no ions form (C12H22O11) • Exception: • Some molecular solutes interact with water to form ions. These would be electrolytes. • Examples: Acids  HCl, H2C3O2 Ammonia  NH3

  8. Strong Electrolytes • Exists in solution completely or almost completely as ions. • All ionic compounds and a few molecular compounds.(Ex: Strong Acids)

  9. Weak Electrolytes • Molecular compounds that produce a small concentration of ions when dissolved in H2O. Ex: Acetic acid (HC2H3O2) only slightly ionizes when dissolved in water. HC2H3O2(aq)H+(aq)+ C2H3O2-(aq) Weak acids are better conductors if they are dilute, as you will see in lab. Explain.

  10. Precipitation Reactions • Reactions that result in an insoluble product. • Insoluble: • Substance with solubility less than 0.01 mol/L • Water molecules cannot overcome the attraction between the ions.

  11. KI (aq) + Pb(NO3)2 (aq)  PbI2(s) + KNO3 (aq) Precipitate

  12. You must be able to determine whether a substance is soluble in water by simple examination of the chemical formula. To do so, you must memorize how specific polyatomic ions act in water. Not as hard as it sounds. We will focus mainly on 10 anions. This will give you the tools to predict the solubility of many compounds.

  13. Solubility of Ionic Compounds • All acetates and nitrates are soluble in water. • All ionic compounds of alkali metals and ammonium are soluble. • (1A goes AWAY) • Solubility rules are on your reference sheet.

  14. Soluble Exceptions Ag+, Hg22+, Pb2+ None Ag+, Hg22+, Pb2+ Ag+, Hg22+, Pb2+ None Sr2+, Ba2+, Hg22+, Pb2+ Cl- Acetate ion C2H3O2- Br- I- NO3- SO42-

  15. Insoluble Exceptions NH4+ and Group 1 metals Group 1 metals and Ba2+, Sr2+, Ca2+ NH4+ and Group 1 metals NH4+ and Group 1 metals and Ba2+, Sr2+, Ca2+ CO32- OH- PO43- S2-

  16. Equation Types • Molecular • Complete Ionic • Net ionic equation 3

  17. Ionic Equations • Those ions that appear on both sides of a complete ionic equation are known as Spectator Ions. • Net ionic equations do not include spectator ions.

  18. Exchange Reactions • Metathesis reactions • Double displacement • Double replacement

  19. Writing Net Ionic Equations • Write a balanced molecular equation. • Rewrite the equation showing ions of strong electrolytes only. • Identify and cancel all spectator ions.

  20. Acid-Base Reactions All this acid rain is killing my complexion! Acids: • Ionize in H2O, causes increase in H+ ions. • H+ ions are bare protons. • Acids are proton donors.

  21. Monoprotic Acids: (HCl, HNO3) • Acids that can only yield one H+ per molecule upon ionization. HCl  H+ + Cl-

  22. Diprotic Acids: (H2SO4) • Ionization occurs in 2 steps. • Only the first ionization is complete.

  23. Is HF a weak or strong acid? weak acid Although it is a weak acid, this acid is extremely reactive because of the F- ion. Must be kept in special polypropylene container because it eats through glass. Used to etch glass. Has caused major accidents in lab.

  24. Bases • Substances that increase the OH- when added to water. (NaOH) • NH3 is a base. In water it accepts an H+ ion from HOH, leaving an OH- in solution. • NH3 is a weak electrolyte • About 1% ionizes to form NH4+/OH-

  25. Strong acids and bases • Acids and bases that ionize completely in solution are strong acids and bases. • Those that only ionize partially are weak acids and bases. • You must memorize these.

  26. Strong Acids Hydrochloric Acid – Hydrobromic Acid – Hydroiodic Acid – Nitric Acid – Sulfuric Acid – Chloric Acid – Perchloric Acid – HCl HBr HI HNO3 H2SO4 HClO3 HClO4

  27. Strong Bases All group 1 Metal Hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) Heavy Group 2 Metal Hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2

  28. Once you memorize the strong acids and bases, you will have enough information to determine if a substance is a strong or weak electrolyte.

  29. Example problems: KF Na3PO4 NH3 CH3CH2OH HCl NO2 HC2H3O2 CH4 NH4Cl CH3Cl strong electrolyte strong electrolyte weak electrolyte nonelectrolyte strong electrolyte nonelectrolyte weak electrolyte nonelectrolyte strong electrolyte nonelectrolyte

  30. Acid + Base Neutralization • Products of a neutralization reaction have none of the properties of an acid or a base. • An acid reacts with a metal hydroxide to form a salt plus water.

  31. Neutralization Reactions • Acid + Base (Metal Hydroxide) Salt + Water • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) • H+ + Cl- + Na+ + OH-  Na+ + Cl- + H2O(l) H+ + OH-  H2O(l)

  32. Write the net ionic equation for the following reaction. It might help to first write the molecular equation, and then the complete ionic equation, followed by the net ionic equation. Potassium Hydroxide + Sulfuric Acid • Ionic equation: • Net Ionic equation:

  33. Neutralization Reaction of Weak Acid *Remember, only strong electrolytes are written as ions.* Acetic Acid + Sodium Hydroxide  HC2H3O2(aq) + NaOH(aq)  NaC2H3O2(aq) + H2O(l) Weak acid strong base soluble salt water HC2H3O2 + Na+ + OH- Na+ + C2H3O2- + H2O(l) HC2H3O2(aq) + OH-(aq)  C2H3O2-(aq)+ H2O(l)

  34. Acid/Base Rx’s with gas formation • Other bases besides OH- react with H+ to form molecular compounds.Two common bases are CO3-2 and S-2. • Carbonates and bicarbonates react with acid to form CO2.

  35. Hydrochloric acid + Sodium Sulfide  2HCl (aq) + Na2S(aq)  H2S(g) + 2NaCl(aq) 2H+ (aq) + S2-(aq)  H2S(g) Hydrochloric acid + Sodium Hydrogen Carbonate  HCl (aq) + NaHCO3(aq)  NaCl(aq) + H2CO3(aq) unstable H2CO3(aq) H2O(l) + CO2(g) Overall HCl (aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g) Net H+ (aq) + HCO3-(aq)  H2O(l) + CO2(g)

  36. Oxidation-Reduction Reactions

  37. Oxidation-Reduction (Redox) Reactions in which electrons are transferred between substances

  38. Use of Oxidation numbers in determining redox reactions is basically a bookkeeping method for keeping track of electrons You must be able to identify an oxidation-reduction reaction. But first, we must learn the rules for assigning oxidation #’s to different species.

  39. Rules for oxidation numbers • Atoms in elemental form are 0. • Monatomic ion; charge of the ion is its oxidation number. • Nonmetals; usually negative numbers. a.) oxygen = -2 unless a peroxide = -1 b.) Hydrogen +1 with nonmetals, -1 with metals c.) Halogens (-1) unless bonded to oxygen (+) in a polyatomic ion (Ex: ClO3-; Cl = +5) 4) Sum of oxidation numbers must = 0 5) Most electronegative (furthest to right and up) element gets a negative charge. See pages 128 – 129 for more on this.

  40. Atoms in elemental form are 0. • Examples • Ag • Pb • Cl2 • O2 • Oxidation # = 0 for 7 diatomic elements and for all • other elements when by themselves.

  41. Monatomic ion-- • charge of the ion is its oxidation number. • Examples • AgCl Ag = +1 Cl = -1 • PbI2 Pb = +2 I = -1 • Fe2O3 Fe = +3 O = -2

  42. 3) Nonmetals; usually negative numbers. a.) oxygen = -2 unless a peroxide = -1 b.) Hydrogen +1 with nonmetals, -1 with metals c.) Halogens (-1) unless bonded to oxygen (+) in a polyatomic ion (Ex: ClO3-; Cl = +5) Examples PbO oxygen = -2 Na2O2 oxygen = -1 H2S hydrogen = +1 NaH hydrogen = -1 KI iodine = -1 KIO2 iodine = + 3

  43. Determine Oxidation # of element red element in each of the following: MnO2 +4 KMnO4 +7 BrO2- +3 BrO3- +5 Br2 0 HClO4 +7 H2SO4 +6 PO33- +3 CaH2 -1 SO42- +6 Na2S -2 Mg(NO3)2 +5

  44. Again, oxidation reduction reactions occur when there is a transfer of electrons from one species to another in a reaction. • If one reactant gains electrons another must lose electrons. • Reduction is always accompanied by oxidation.

  45. Oxidation-Reduction Reactions • An atom that becomes more positively charged is oxidized. • This is due to loss of e-. • The gain of electrons by an atom is called reduction.

  46. Two mnemonics for remembering which substance is undergoing oxidation and which is undergoing reduction? - OIL -- RIG Oxidation Involves Loss -- Reduction Involves Gain “Leo the lion says Ger” Loss of electrons oxidation -- Gain of electrons reduction

  47. Many metals react with O2 in • the air to form metal oxides. • Metals lose electrons to oxygen. • 2 Fe + O2 2 FeO • As Fe is oxidized (loses e-), • oxygen is reduced (gains e-). • Reduction is gain

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